contact process | chemistry | Britannica.com Contact process varnish Contact process, modern industrial method of producing sulfuric acid; it has largely replaced the chamber, or lead-chamber, process. Sulfur dioxide and oxygen , passed over a hot catalyst , unite to form sulfur trioxide , which in turn combines with water to make sulfuric acid. Contact-process plants are of two types. The simpler type, the sulfur -burning contact plants, use sulfur as the raw material. Molten sulfur is burned to form sulfur dioxide, which is cooled, then oxidized, usually in the presence of pellets of porous siliceous material impregnated with vanadium pentoxide and a potassium compound , to form sulfur trioxide at moderately high temperatures. The other type of contact-process plant produces sulfur dioxide from low-grade, sulfur-bearing materials, such as pyrite. Cooling of the gas is necessary to remove impurities and to condense and remove part of the water vapour, which would dilute the acid product. The sulfur dioxide gas is then dried with concentrated sulfuric acid.

As a result of its purification, the gas in this process is cold, instead of hot as in sulfur-burning plants, and must be heated to the temperature at which the catalyst begins to function. Learn More in these related articles: sulfuric acid dense, colourless, oily, corrosive liquid; one of the most important of all chemicals, prepared industrially by the reaction of water with sulfur trioxide (see sulfur oxide), which in turn is made by chemical combination of sulfur dioxide and oxygen either by the contact process or the chamber... oxygen (O) nonmetallic chemical element of Group 16 (VIa, or the oxygen group) of the periodic table. Oxygen is a colourless, odourless, tasteless gas essential to living organisms, being taken up by animals, which convert it to carbon dioxide; plants, in turn, utilize carbon dioxide as a source of carbon and... 1 Reference found in Britannica Articles Assorted Reference sulfuric acid production (in chemical industry: Contact process ) External Links Corrections? Updates? Help us improve this article! Contact our editors with your feedback. MEDIA FOR: You have successfully emailed this. Error when sending the email. Try again later.

Edit Mode Submit Tips For Editing We welcome suggested improvements to any of our articles. You can make it easier for us to review and, hopefully, publish your contribution by keeping a few points in mind. Encyclopædia Britannica articles are written in a neutral objective tone for a general audience. You may find it helpful to search within the site to see how similar or related subjects are covered. Any text you add should be original, not copied from other sources. At the bottom of the article, feel free to list any sources that support your changes, so that we can fully understand their context. (Internet URLs are the best.) Your contribution may be further edited by our staff, and its publication is subject to our final approval. Unfortunately, our editorial approach may not be able to accommodate all contributions. Submit Thank You for Your Contribution! Our editors will review what you've submitted, and if it meets our criteria, we'll add it to the article. Please note that our editors may make some formatting changes or correct spelling or grammatical errors, and may also contact you if any clarifications are needed. Uh Oh There was a problem with your submission. Please try again later.

Close Date Published: July 20, 1998 URL: Access Date: January 17, 2017 Share Sulphuric acid- Contact process- industrial preparation of sulphuric acid-H2SO4 DILUTION OF OLEUM. For latest information , free computer courses and high impact notes visit : PREPARATION OF SO2 SO2 is obtained by burning sulphur or by heating iron pyrite (FeS2) in pyrite burner. S + O2 èSO2 4FeS2 + 11O2 è2Fe2O3 + 8SO2 PURIFICATION OF SO2 SO2 contains a number of impurities such as dust particles, Arsenous oxide, vapours, sulphur etc. These impurities must be removed otherwise catalyst loses its efficiency (catalyst poisoning). DUST CHAMBER: SO2 is first passed through the dust chamber where steam is spread over the gas to remove dust particles, which settle down. Fe(OH)3 also sprayed over to remove oxides of Arsenic. WASHING TOWER: SO2 is then passed through a washing tower after cooling.

Here it is sprayed by water to remove any other soluble impurities. DRYING TOWER: The gas is now dried by passing through drying tower where conc. H2SO4 (dehydrating agent) is sprayed. H2SO4 removes moisture from SO2. TYNDALL BOX (TEST BOX): Arsenic oxide is a poison for the catalyst. It is removed when the gas is passed over ferric hydroxide. As2O3 + 2Fe(OH)3 è 2FeAsO3 + 3H2O. In order to remove traces of As2O3, it is passed through a test box, where a strong beam of light is thrown against the gas. If there is no scattering of light in the box, it indicates that gas is free from As2O3. OXIDATION OF SO2 TO SO3 CONTACT TOWER: Oxidation of SO2 is carried out in contact tower where V2O5 is filled in different pipes. SO2 here reacts with air (O2) to produce SO3. Under above conditions 98% SO2 is converted into SO3.

2SO2 + O2 è 2SO3 + 45Kcal CONDITIONS NECESSARY FOR MAXIMUM YIELD OF SO3: Oxidation of SO2 is a reversible and exothermic process in which volume of product is less than the volumes of reactants. In order to obtain maximum amount of SO3, according to Le-Chatelierï¿½s Principle following conditions are necessary. CONCENTRATION: Excess of O2. TEMPEATURE: A decrease in temperature favours reaction in forward direction. Optimum temperature for this process is 450oC to 500oC. PRESSURE: Since volumes of reactants are greater than the product (3:2), therefore, according to Le-Chatelierï¿½s Principle a high pressure is favourable. Optimum pressure is about 1.5 to 1.7 atmosphere. USE OF CATALYST: At low temperature, rate of reaction decreases. To increase rate of reaction a catalyst vanadium pentaoxide (V2O5) is used. ABSORPTION OF SO3 IN H2SO4 SO3 is not directly passed in water, because a dense fog of minute particles of H2SO4 is produced.

It is therefore, dissolved in conc.H2SO4 to form pyrosulphuric acid (oleum). SO3 + H2SO4 è H2S2O7 (OLEUM) DILUTION OF OLEUM Oleum is now diluted with water to form H2SO4 of required concentration. H2S2O7 + H2O è 2H2SO4 For latest information , free computer courses and high impact notes visit : lead: Uses lead Uses The single most important commercial use of lead is in the manufacture of lead-acid storage batteries (see battery, electric ). It is also used in alloys such as fusible metals, antifriction metals, solder , and type metal . Shot lead is an alloy of lead, antimony, and arsenic. Lead foil is made with lead alloys. Lead is used for covering cables and as a lining for laboratory sinks, tanks, and the "chambers" in the lead-chamber process for the manufacture of sulfuric acid. It is used extensively in plumbing . Because it has excellent vibration-dampening characteristics, lead is often used to support heavy machinery and was used in the foundations of the Pan Am Building built over Grand Central Station in New York City.

Lead is also employed as protective shielding against X rays and radiation from nuclear reactors. Lead has many commonly used compounds. Commercially important are the lead oxides, which have many uses. Litharge is lead monoxide, PbO; red lead is lead tetroxide, Pb3O4; lead peroxide or dioxide, PbO2, is used in matches, as a mordant in dyeing, and as an oxidizing agent. White lead , 2PbCO3·Pb(OH)2 (basic lead carbonate), is an important pigment used in paints, putty, and ceramics. Chrome yellow, PbCrO4, is a bright yellow pigment. "Sublimed white lead," PbSO4·Pb(OH)2 (basic lead sulfate), is also used as a pigment. Lead acetate (sugar of lead) is used as a mordant, and lead azide, Pb(N3)2, is employed as a detonator for explosives. Lead arsenate is used as an insecticide. Tetraethyl lead, used as a antiknock compound in gasoline, is now banned for environmental reasons in the United States and other countries.

Although lead and most of its compounds are only slightly soluble in water, the use of lead pipe to carry drinking water is dangerous, since lead is a cumulative poison that is not excreted from the body (see lead poisoning ). The "lead" of lead pencils does not contain lead; it is a mixture of graphite and clay. Sections in this article: SULFURIC ACID (H2SO4) 1830, UN 1831 (fuming), UN 1832 (spent) GENERAL DESCRIPTION OF SULFURIC ACID & ITS SALTS Sulfuric acid , also spelled Sulphuric acid, sometimes called Oil Of Vitriol, or Hydrogen Sulfate battery acid, is a dense, colourless, oily, corrosive liquid; more sulfuric acid is produced than any other chemical. It has widely varied uses and plays some part in the production of nearly all manufactured goods. Pure sulfuric acid with a specific gravity of 1.85. freezes at 10.37° C , boils at 340° C and is soluble in all proportions in water.

When heated, the pure acid partially decomposes into water and sulfur trioxide. Sulfuric acid is a very strong acid in aqueous solution; it is largely changed to hydrogen ions (H+) and sulfate ions (SO42-). Each molecule gives two H+ ions, thus sulfuric acid is dibasic, which forms both normal sulfates (with both hydrogens replaced, e.g., sodium sulfate, Na2SO4) and acid sulfates, also called bisulfates or hydrogen sulfates (with only one hydrogen replaced, e.g., sodium bisulfate, NaHSO4). Dilute solutions of sulfuric acid show all the behavior characteristics of acids. It turns blue litmus red. It conducts electricity, neutralizes alkalies, corrodes active many metals, releasing hydrogen gas, and forming the sulfates. It reacts with most hydroxides and oxides, with some carbonates and sulfides, and with some salts. Concentrated sulfuric acid, formerly called oil of vitriol, is a weak acid and a poor electrolyte because relatively little of it is dissociated into ions.

When concentrated sulfuric acid is heated, it behaves also as an oxidizing agent dissolving relatively unreactive metals as copper, mercury, and lead to produce metal sulfate, sulfur dioxide, and water. Because the concentrated acid has a fairly high boiling point, it can be used to release more volatile acids from their salts, or common salts, when are heated with concentrated sulfuric acid, HCl gas is evolved. As concentrated sulfuric acid has a very strong affinity for water, it is a valuable desiccating agent, is used as a drying agent and can be used to dehydrate many compounds. It removes water from, and therefore chars, wood, cotton, sugar, and paper. It is used in the manufacture of ether, nitroglycerin, and dyes for its property as a desiccant. Sulfur trioxide dissolves readily in concentrated sulfuric acid to form pyrosulfuric acid, H2S2O7, which is also called fuming sulfuric acid or oleum . Sulfuric acid is prepared industrially by the reaction of water with sulfur trioxide, which in turn is made by chemical combination of sulfur dioxide and oxygen either by the contact process or the chamber process.

The lead chamber process is used to produce much of the acid used to make fertilizers. It produces a relatively dilute acid (62% - 78%). The contact process produces a more concentrated acid but requires purer raw materials and the use of expensive catalysts. Some sulfuric acid is also made from ferrous sulfate waste solutions from pickling iron and steel and from waste acid sludge from oil refineries. The uses of sulfuric acid are so varied that the volume of its production provides an approximate index of general industrial activity. Its main use is in phosphate fertilizer production, both superphosphate of lime and ammonium sulfate. It is widely also used to manufacture chemicals, e.g., in making hydrochloric acid, nitric acid, sulfate salts, synthetic detergents, dyes and pigments, explosives, drugs, other acids, parchment paper, glue and wood preservatives. It is used in the purification of petroleum to wash impurities out of gasoline and other refinery products.Sulfuric acid is used in processing metals, e.g., in pickling (cleaning) of metal, electroplating baths, nonferrous metallurgy. Rayon is made with sulfuric acid.

In one of its most familiar applications, it serves as the electrolyte in the lead-acid storage battery commonly used in motor vehicles (acid for this use, containing about 33% H2SO4 and with specific gravity about 1.25, is often called battery acid). Any of numerous chemical compounds related to sulfuric acid, formed by replacing one or both of the hydrogens with a metal or a radical are called sulfates ( also spelled sulphates). Sulfates are salts or esters of sulfuric acid. One group of these derivatives is composed of salts containing the sulfate ion, and positively charged ions such as those of sodium, magnesium, or ammonium; a second group is composed of esters, in which the hydrogen atoms of sulfuric acid have been replaced by carbon-containing combining groups such as methyl (CH3) or ethyl (C2H5). Most metal sulfates are readily soluble in water, but calcium and mercuric sulfates are only slightly soluble, while barium, lead, strontium, and mercurous sulfates are insoluble.

In chemical analysis, the sulfate ion, SO42-, is usually detected by adding barium chloride solution; the white barium sulfate precipitate that forms is insoluble in hydrochloric acid. Sulfates are widely distributed in nature. Barium sulfate occurs as barite; calcium sulfate is found as gypsum, alabaster, and selenite; Epsom salts is magnesium sulfate; sodium sulfate occurs as its decahydrate, Glauber salt; and strontium sulfate occurs as celestite. Some sulfates were formerly known as vitriols; blue vitriol is cupric sulfate, green vitriol is ferrous sulfate, and white vitriol is zinc sulfate. Alums are solid sulfates, containing two different metals and two sulfate radicals. Organic sulfates are esters. They can be formed by reacting an alcohol with cold sulfuric acid. They are also formed by the reaction of sulfuric acid with a solid bond in an alkene; the product is called an alkyl hydrogen sulfate. An alkyl hydrogen sulfate can be broken down to an alcohol and sulfuric acid by heating it with water (hydrolysis); this reaction is often used to synthesize alcohols.

sulfuric acid facts, information, pictures | Encyclopedia.com articles about sulfuric acid The Columbia Encyclopedia, 6th ed. Copyright The Columbia University Press sulfuric acid, chemical compound, H2SO4, colorless, odorless, extremely corrosive, oily liquid. It is sometimes called oil of vitriol. Concentrated Sulfuric Acid When heated, the pure 100% acid loses sulfur trioxide gas, SO3, until a constant-boiling solution, or azeotrope, containing about 98.5% H2SO4 is formed at 337°C. Concentrated sulfuric acid is a weak acid (see acids and bases ) and a poor electrolyte because relatively little of it is dissociated into ions at room temperature. When cold it does not react readily with such common metals as iron or copper. When hot it is an oxidizing agent, the sulfur in it being reduced; sulfur dioxide gas may be released. Hot concentrated sulfuric acid reacts with most metals and with several nonmetals, e.g., sulfur and carbon.

Because the concentrated acid has a fairly high boiling point, it can be used to release more volatile acids from their salts, e.g., when sodium chloride (NaCl), or common salt, is heated with concentrated sulfuric acid, hydrogen chloride gas, HCl, is evolved. Concentrated sulfuric acid has a very strong affinity for water. It is sometimes used as a drying agent and can be used to dehydrate (chemically remove water from) many compounds, e.g., carbohydrates. It reacts with the sugar sucrose, C12H22O11, removing eleven molecules of water, H2O, from each molecule of sucrose and leaving a brittle spongy black mass of carbon and diluted sulfuric acid. The acid reacts similarly with skin, cellulose, and other plant and animal matter. When the concentrated acid mixes with water, large amounts of heat are released; enough heat can be released at once to boil the water and spatter the acid. To dilute the acid, the acid should be added slowly to cold water with constant stirring to limit the buildup of heat. Sulfuric acid reacts with water to form hydrates with distinct properties.

Dilute Sulfuric Acid Dilute sulfuric acid is a strong acid and a good electrolyte; it is highly ionized, much of the heat released in dilution coming from hydration of the hydrogen ions . The dilute acid has most of the properties of common strong acids. It turns blue litmus red. It reacts with many metals (e.g., with zinc), releasing hydrogen gas, H2, and forming the sulfate of the metal. It reacts with most hydroxides and oxides, with some carbonates and sulfides, and with some salts. Since it is dibasic (i.e., it has two replaceable hydrogen atoms in each molecule), it forms both normal sulfates (with both hydrogens replaced, e.g., sodium sulfate, Na2SO4) and acid sulfates, also called bisulfates or hydrogen sulfates (with only one hydrogen replaced, e.g., sodium bisulfate, NaHSO4). Production of Sulfuric Acid There are two major processes (lead chamber and contact) for production of sulfuric acid, and it is available commercially in a number of grades and concentrations.

The lead chamber process, the older of the two processes, is used to produce much of the acid used to make fertilizers; it produces a relatively dilute acid (62%–78% H2SO4). The contact process produces a purer, more concentrated acid but requires purer raw materials and the use of expensive catalysts. In both processes sulfur dioxide is oxidized and dissolved in water. The sulfur dioxide is obtained by burning sulfur, by burning pyrites (iron sulfides), by roasting nonferrous sulfide ores preparatory to smelting, or by burning hydrogen sulfide gas. Some sulfuric acid is also made from ferrous sulfate waste solutions from pickling iron and steel and from waste acid sludge from oil refineries.

Lean Chamber Process In the lead chamber process hot sulfur dioxide gas enters the bottom of a reactor called a Glover tower where it is washed with nitrous vitriol (sulfuric acid with nitric oxide, NO, and nitrogen dioxide, NO2, dissolved in it) and mixed with nitric oxide and nitrogen dioxide gases; some of the sulfur dioxide is oxidized to sulfur trioxide and dissolved in the acid wash to form tower acid or Glover acid (about 78% H2SO4). From the Glover tower a mixture of gases (including sulfur dioxide and trioxide, nitrogen oxides, nitrogen, oxygen, and steam) is transferred to a lead-lined chamber where it is reacted with more water. The chamber may be a large, boxlike room or an enclosure in the form of a truncated cone. Sulfuric acid is formed by a complex series of reactions; it condenses on the walls and collects on the floor of the chamber. There may be from three to twelve chambers in a series; the gases pass through each in succession. The acid produced in the chambers, often called chamber acid or fertilizer acid, contains 62% to 68% H2SO4.

After the gases have passed through the chambers they are passed into a reactor called the Gay-Lussac tower where they are washed with cooled concentrated acid (from the Glover tower); the nitrogen oxides and unreacted sulfur dioxide dissolve in the acid to form the nitrous vitriol used in the Glover tower. Remaining waste gases are usually discharged into the atmosphere. Contact Process In the contact process, purified sulfur dioxide and air are mixed, heated to about 450°C, and passed over a catalyst; the sulfur dioxide is oxidized to sulfur trioxide. The catalyst is usually platinum on a silica or asbestos carrier or vanadium pentoxide on a silica carrier. The sulfur trioxide is cooled and passed through two towers. In the first tower it is washed with oleum (fuming sulfuric acid, 100% sulfuric acid with sulfur trioxide dissolved in it). In the second tower it is washed with 97% sulfuric acid; 98% sulfuric acid is usually produced in this tower. Waste gases are usually discharged into the atmosphere. Acid of any desired concentration may be produced by mixing or diluting the products of this process.

Uses of Sulfuric Acid Sulfuric acid is one of the most important industrial chemicals. More of it is made each year than is made of any other manufactured chemical; more than 40 million tons of it were produced in the United States in 1990. It has widely varied uses and plays some part in the production of nearly all manufactured goods. The major use of sulfuric acid is in the production of fertilizers, e.g., superphosphate of lime and ammonium sulfate. It is widely used in the manufacture of chemicals, e.g., in making hydrochloric acid, nitric acid, sulfate salts, synthetic detergents, dyes and pigments, explosives, and drugs. It is used in petroleum refining to wash impurities out of gasoline and other refinery products. Sulfuric acid is used in processing metals, e.g., in pickling (cleaning) iron and steel before plating them with tin or zinc. Rayon is made with sulfuric acid.

It serves as the electrolyte in the lead-acid storage battery commonly used in motor vehicles (acid for this use, containing about 33% H2SO4 and with specific gravity about 1.25, is often called battery acid). History of Sulfuric Acid Although sulfuric acid is now one of the most widely used chemicals, it was probably little known before the 16th cent. It was prepared by Johann Van Helmont (c.1600) by destructive distillation of green vitriol (ferrous sulfate) and by burning sulfur. The first major industrial demand for sulfuric acid was the Leblanc process for making sodium carbonate (developed c.1790). Sulfuric acid was produced at Nordhausen from green vitriol but was expensive. A process for its synthesis by burning sulfur with saltpeter (potassium nitrate) was first used by Johann Glauber in the 17th cent. and developed commercially by Joshua Ward in England c.1740. It was soon superseded by the lead chamber process, invented by John Roebuck in 1746 and since improved by many others.

The contact process was originally developed c.1830 by Peregrine Phillips in England; it was little used until a need for concentrated acid arose, particularly for the manufacture of synthetic organic dyes. Cite this article Pick a style below, and copy the text for your bibliography. MLA The Contact Process for the manufacture of sulphuric acid This can either be made by burning sulphur in an excess of air: . . . or by heating sulphide ores like pyrite in an excess of air: In either case, an excess of air is used so that the sulphur dioxide produced is already mixed with oxygen for the next stage. Converting the sulphur dioxide into sulphur trioxide This is a reversible reaction, and the formation of the sulphur trioxide is exothermic. A flow scheme for this part of the process looks like this: The reasons for all these conditions will be explored in detail further down the page. Converting the sulphur trioxide into sulphuric acid This can't be done by simply adding water to the sulphur trioxide - the reaction is so uncontrollable that it creates a fog of sulphuric acid.

Instead, the sulphur trioxide is first dissolved in concentrated sulphuric acid: The product is known as fuming sulphuric acid or oleum. This can then be reacted safely with water to produce concentrated sulphuric acid - twice as much as you originally used to make the fuming sulphuric acid. Explaining the conditions The proportions of sulphur dioxide and oxygen The mixture of sulphur dioxide and oxygen going into the reactor is in equal proportions by volume. Avogadro's Law says that equal volumes of gases at the same temperature and pressure contain equal numbers of molecules. That means that the gases are going into the reactor in the ratio of 1 molecule of sulphur dioxide to 1 of oxygen. That is an excess of oxygen relative to the proportions demanded by the equation. According to Le Chatelier's Principle, Increasing the concentration of oxygen in the mixture causes the position of equilibrium to shift towards the right. Since the oxygen comes from the air, this is a very cheap way of increasing the conversion of sulphur dioxide into sulphur trioxide. Why not use an even higher proportion of oxygen? This is easy to see if you take an extreme case. Suppose you have a million molecules of oxygen to every molecule of sulphur dioxide.

The equilibrium is going to be tipped very strongly towards sulphur trioxide - virtually every molecule of sulphur dioxide will be converted into sulphur trioxide. Great! But you aren't going to produce much sulphur trioxide every day. The vast majority of what you are passing over the catalyst is oxygen which has nothing to react with. By increasing the proportion of oxygen you can increase the percentage of the sulphur dioxide converted, but at the same time decrease the total amount of sulphur trioxide made each day. The 1 : 1 mixture turns out to give you the best possible overall yield of sulphur trioxide. The temperature Equilibrium considerations You need to shift the position of the equilibrium as far as possible to the right in order to produce the maximum possible amount of sulphur trioxide in the equilibrium mixture. The forward reaction (the production of sulphur trioxide) is exothermic. According to Le Chatelier's Principle, this will be favoured if you lower the temperature. The system will respond by moving the position of equilibrium to counteract this - in other words by producing more heat. In order to get as much sulphur trioxide as possible in the equilibrium mixture, you need as low a temperature as possible.

However, 400 - 450°C isn't a low temperature! Rate considerations The lower the temperature you use, the slower the reaction becomes. A manufacturer is trying to produce as much sulphur trioxide as possible per day. It makes no sense to try to achieve an equilibrium mixture which contains a very high proportion of sulphur trioxide if it takes several years for the reaction to reach that equilibrium. You need the gases to reach equilibrium within the very short time that they will be in contact with the catalyst in the reactor. The compromise 400 - 450°C is a compromise temperature producing a fairly high proportion of sulphur trioxide in the equilibrium mixture, but in a very short time. The pressure Equilibrium considerations Notice that there are 3 molecules on the left-hand side of the equation, but only 2 on the right. According to Le Chatelier's Principle, if you increase the pressure the system will respond by favouring the reaction which produces fewer molecules. That will cause the pressure to fall again. In order to get as much sulphur trioxide as possible in the equilibrium mixture, you need as high a pressure as possible. High pressures also increase the rate of the reaction.

However, the reaction is done at pressures close to atmospheric pressure! Economic considerations Even at these relatively low pressures, there is a 99.5% conversion of sulphur dioxide into sulphur trioxide. The very small improvement that you could achieve by increasing the pressure isn't worth the expense of producing those high pressures. The catalyst Equilibrium considerations The catalyst has no effect whatsoever on the position of the equilibrium. Adding a catalyst doesn't produce any greater percentage of sulphur trioxide in the equilibrium mixture. Its only function is to speed up the reaction. Rate considerations In the absence of a catalyst the reaction is so slow that virtually no reaction happens in any sensible time. The catalyst ensures that the reaction is fast enough for a dynamic equilibrium to be set up within the very short time that the gases are actually in the reactor. Note: If you are interested in the mechanism for the catalytic reaction you will find it on an introductory page on types of catalysts. Use the BACK button on your browser if you want to return to this page. Questions to test your understanding If this is the first set of questions you have done, please read the introductory page before you start. You will need to use the BACK BUTTON on your browser to come back here afterwards.

Sulphuric Acid Sulphuric Acid A molecule of Sulphuric Acid, H2SO4, consists of two atoms of hydrogen, one atom of sulphur and four atoms of oxygen Sulphuric acid is a colourless viscous corrosive oily liquid, which has Melting Point : 10.3 degC Specific gravity or density 1.94 Flash point none Sulphuric acid is the strong acid produced by dissolving sulphur trioxide in water. SO3 + H2O ==> H2SO4 The Strength of Acids is determined by the degree to which they are ionised in aqueous solution. For example, Sulphuric Acid, H2SO4, which is a strong acid is fully dissociated, and all the displaceable hydrogen in the acid is present in solution as Hydrogen Ion, H(+). H2SO4 ==> H(+) + SO4 100% as H(+) In contrast, the weak acids ethanoic acid, CH3COOH, is partially ionised in solution, and only approximately 5% of the displaceable Hydrogen in the acid is present in solution as hydrogen ion, H(+).

CH3COOH ==> H(+) + CH3COO(-) 5% as H(+) Sulphuric acid is a powerful protonating agent. It is also a moderately strong oxidizing agent . Sulphuric acid is also a powerful dehydrating agent and is used to remove a molecule of water from many organic compounds. In dilute solution, sulphuric acid is a strong dibasic acid forming two series of salts. A Dibasic Acid has two acidic hydrogen atoms in its molecules which can be ionised. Sulphuric Acid, H2SO4, is a dibasic acid, because it contains two hydrogens atoms which ionise in aqueous solution to become Hydrogen Ions, H(+). H2SO4 ==> 2 H(+) + SO4(2 -) Sulphuric acid is an important industrial chemical and it has many uses as a strong oxidising agent and a powerful dehydrating agent. Commercially available sulphuric acid is as a 96-98% solution of the acid in water. It is a powerful protonating agent.

It is also a powerful dehydrating agent and is used to remove a molecule of Water, HO2, from many organic compounds. The Dehydration Reactions of Alcohols results in their converted into an alkene, and involves the elimination of a molecule of water. Dehydration requires the presence of an acid and the application of heat. Preparation of Sulphuric Acid Combustion of Sulphur When a small amount of Sulphur, S, is kindled on a deflagrating spoon, it burns with a bright blue flame when introduced into a gas jar containing Oxygen, O2. A gas, Sulphur Dioxide, SO2, is the main product of the combustion. However, a little Sulphur Trioxide, SO3, is also formed, which makes the gas slightly cloudy. S + O2 ==> SO2 Sulphur Dioxide 2S + 3O2 ==> 2SO3 Sulphur Trioxide When shaken with water, the products of combustion dissolve, forming an acidic solution which turns litmus red.

SO2 + H2O ==> H2SO3 Sulphur Sulphurous Dioxide Acid SO2 + H2O ==> H2SO4 Sulphur Sulphuric Trioxide Acid Manufacture of Sulphuric Acid Sulphuric acid was manufactured by the lead-chamber process until the mid-1930s, but this process has now been replaced by the contact process, involving the catalytic oxidation of sulphur dioxide. Properties of Sulphuric Acid The Contact Process is used for manufacturing sulphuric acid and fuming sulphuric acid from sulphur dioxide, which is made by burning sulphur or by roasting sulphide ores and oxygen (in the form of air) which combine to form sulphur trioxide in the presence of a catalyst . The reaction is exothermic and the conditions are controlled to keep the temperature at 450 degC. The catalyst used is valadium oxide (V2O5). The sulphur trioxide is dissolved in sulphuric acid to form fuming sulphuric acid, this is called oleum.

S(s) + O2(g) ==> SO2(g) Sulphur Dioxide 2SO2(g) + O2(g) ==> 2SO3(g) Sulphur Trioxide SO3(g) + H2SO4(l)==> H2S2O7(l) Oleum This Oleum, H2S2O7, is then diluted with Water, H2O, to produce concentrated Sulphuric Acid, H2SO4. H2S2O7(l)+ H2O (l) ==> 2 H2SO4(l) Oleum Sulphuric Acid Reactions of Sulphuric acid Electrolysis of a Solution of dilute Sulphuric Acid The Electrolysis of an Aqueous Solution of dilute Sulphuric Acid is often carried out in a Hofmann Voltammeter , an apparatus in which the gases evolved at the anode and cathode can be collected in separate graduated tubes. When the solution is electrolyzed hydrogen is produced at the cathode and oxygen at the anode. These gases can be shown to be present in a 2 to 1 ratio and result from the electrolysis of water under acidic conditions.

Sulphuric acid is a strong electrolyte is fully dissociated in aqueous solution. H2SO4 ==> 2 H(+) + SO4(2 -) Water is a weak electrolyte and is only slightly dissociated H2O ==> H(+) + OH(-) During electrolysis, the Hydrogen Ions, H(+), migrates towards the cathode, and are discharged there (i.e. they gain an electron and are converted to hydrogen gas). 2 H(+) + 2 e(-) ==> H2- At the anode the concentration of Hydroxyl Ions, HO(-),is too low to maintain a reaction and the Sulphate Ions, SO4(2 -) are not oxidized but remain on in solution at the end. Water molecules must be the species reacting at the anode.

2 H2O ==> O2 + 4 H(+) + 4 e(-) The overall reaction is 2 H(+) + 2e(-) ==> H2 4 H(+) + 4e(-) ==> 2H2 Anode Reaction : 2 H2O ==> O2 + 4 H(+) + 4 e(-) Overall Cell Reaction: 4 H(+) + 2 H2O ==> 2 H2 + O2 + 4 H(+) For every Hydrogen Ions, H(+), discharged at the anode, another hydrogen ion is formed at the cathode. The net result is that the concentration of the Sulphuric Acid, H2SO4, remains constant and this electrolysis consists of the decomposition of water with the overall reaction 2H2O ==> 2H2- + O2- Ferrous Sulphate, Fe(II)SO4, is the salt formed when Iron, Fe, is dissolved in Sulphuric Acid, H2SO4.

Hydrogen Chloride, HCl, may be prepared in the laboratory by heating Concentrated Sulphuric Acid, H2SO4, with Sodium Chloride, NaCl. NaCl + H2SO4 ==> NaHSO4 + HCl Many Metallic Chlorides liberate Chlorine, Cl2, when treated with Sulphuric Acid, H2SO4, and Manganese Dioxide, MnO2). Many Metallic Chlorides liberate Hydrogen Chloride gas, HCl, when warmed with concentrated Sulphuric Acid, H2SO4. Sulphur Trioxide, SO3, is prepared by heating concentrated Sulphuric Acid, H2SO4, with a large excess of Phosphorus Pentoxide, P2O5. H2SO4 + P2O5 ==> SO3 + 2 HPO3 Sulphur Dioxide, SO2, is usually made in the laboratory by heating concentrated Sulphuric Acid, H2SO4, with Copper turnings, Cu.

Cu + 2 H2SO4 ==> CuSO4 + SO2 + 2 H2O Hydrogen Fluoride, HF, can be prepared in the laboratory by heating Concentrated Sulphuric Acid, H2SO4, with Calcium Fluoride, CaF2. H2SO4 + CaF2 ==> 2 HF + CaSO4 Hydrogen Iodide, HI, can be prepared by direct combination of the elements using a platinum catalyst. In the laboratory it is prepared by heating Concentrated Sulphuric Acid, H2SO4, with Sodium Iodide, NaI. H2SO4 + 2 NaI ==> 2 HI + Na2SO4 Methanol, CH3OH, does not undergo dehydration reactions. Instead, in reaction with Sulphuric Acid, H2SO4, the ester, Dimethyl Sulphate, (CH3)2SO4, is formed.

concentrated H2SO4 2CH3OH ==> (CH3)2SO4 + H2O Methanol Dimethyl Water Sulphate Sulphuric Acid, H2SO4, absorbs Ethylene, C2H4, at room temperature to form Ethyl Hydrogen Sulphate, C2H5.HSO4, with much evolution of heat. C2H4 + H2SO4 ==> C2H5.HSO4 If this is treated with Water, H2O and warmed, Ethanol, C2H5OH, is formed. heat C2H5.HSO4 + H2O ==> C2H5OH + H2SO4 Uses of Sulphuric Acid The Daniell Cell, which is a primary voltaic cell having a positive electrode of Copper, Cu, and a negative electrode of Zinc Amalgam, Zn (in alloy with Hg), was invented by the British chemist John Daniel in 1836AD. The Zinc Amalgam electrode is placed in an electrolyte of dilute Sulphuric Acid solution, H2SO4, or Zinc Sulphate solution, ZnSO4, in a porous earthenware pot.

This porous pot stands in a solution of Copper Sulphate, CuSO4, in which the Copper electrode is immersed. The Zinc electrode, Zn, acts as a source of electrons, which flow through an external wire which connects the two electrodes, while the Zinc Ions, Zn(2 +), from the electrode go into solution. Cathode reaction : Zn ==> Zn(++) + 2 e(-) On reaching the Copper Electrode, these electrons combine with Copper Ions, Cu(2 +), in solution and the discharged copper ions are deposited on the copper electrode as Copper metal, Cu. Anode reaction : Cu(++) + 2 e(-) ==> Cu An equation for the overall chemical process is obtained by adding together the two half-cell reactions in such a way that the electrons "cancel out". Zn + Cu(++) ==> Zn(++) + Cu While the reaction takes place ions move through the porous pot, but when it is not in use the cell should be dismantled to prevent the diffusion of one electrolyte into the other.

Sulphuric acid Sulphuric acid Sulphuric acid Sulphuric acid is probably the most important of all chemicals, because of its extensive use in a very large number of manufacturing operations. Of the immense quantities made yearly, the greater part does not come upon the market; for, being expensive and difficult to ship, consumers of large amounts generally make their own acid. The commercial grades of acid have special names. A moderately strong acid (50°_55° He.), such as condenses in the lead chambers, is known as "chamber acid." It contains from 62 to 70 per cent of H2S04, and is strong enough for use in the manufacture of fertilizer, and for other purposes requiring a dilute acid. By concentrating this chamber acid, an acid of 60° Be. is obtained, containing about 78 per cent of H2S04> which is sufficiently strong for most technical uses. Further evaporation in platinum or iron pans yields an acid of 66° Be., containing 93.5 per cent of H2S04, and known as oil of vitriol.

Faming or Nordhausen acid, which is still more concentrated, is prepared by special means. It is essentially a solution of sulphuric anhydride (S03) in sulphuric acid. This is the acid which was prepared by the alchemists in the Middle Ages. In about the year 1740, Ward, an Englishman, began to make sulphuric acid on a moderately large scale. He burned sulphur and nitre (KNO3) together, and condensed the vapors in glass vessels containing a little water. The dilute acid so formed was then concentrated in glass alembics or retorts. In this way all acid was produced at a lower price than the fuming acid could be made. and the industry was soon established on [t commercial scale. The reactions involved ill Ward's process are the bases of the method now in use; this consists in bringing together, under suitable conditions, sulphur dioxide, oxygen, and water as steam, in the presence of certain oxides of nitrogen. The latter probably act as carriers of the oxygen, causing it to unite with the sulphur dioxide and water to form the acid.

The apparent reaction is;- S02 + H20 + 0 = H2S04- Hut this does not represent the actual process, which is more complicated than it at first appears. Several theories have been advanced to explain the reactions occurring in the lead chambers, and the part taken by the nitrogen oxides, but the most generally accepted one, that of Lunge, regards nitrous anhydride (N2O3) * as the essential factor. t According to this view, the principal reactions involved are as follows: 1) 2S02 + N2O3 + O2 + H20 = 2S02.(OH).(ONO) (Nitrosylsulphuric acid); 2) 2 S02.(OH).(ONO) + H20 = 2 S02(OH)2 + N2O3; or, 3) 2 SO2(OH)(ONO) + S02 +0 +H20 = 3 S02(OH)2 +N203.

First there is a union of sulphur dioxide, nitrous anhydride, oxygen, and water, to form nitrosylsulphuric acid, which probably separates as part of the mist or fog seen in the lead chambers. But in the presence of water vapor or of dilute sulphuric acid, this nitrosylsulphuric acid is at once decomposed, according to reaction (2), sulphuric acid being formed, and nitrous anhydride regenerated; or if sulphur dioxide and oxygen are concerned in the process, then reaction (3) occurs. This cycle of reactions repeats an indefinite number of times. But in the first lead chamber, where the temperature is rather high and an excess of water vapor is usually present, the following secondary reactions probably occur to a greater or less extent:- 4) 2 S02.(OH).(ONO) + S02 + 2 H20 = 3 H2S04 + 2 NO, this reaction being only momentary.

Since there is usually an excess of oxygen present, however, the nitric oxide here formed is at once brought into action again, thus:- 5) 2 S02 + 2 NO + 30 + H20 = 2 S02.(OH).(0N0). If there is a deficiency of oxygen, the nitric oxide is not returned to the process, but passes through the several chambers and, since it is not absorbed by the concentrated acid in the Gay-Lussac tower, it escapes into the atmosphere and is lost. The nitrogen oxides are derived from nitric acid, or by the action of sulphuric acid on sodium nitrate in the nitre pots. When nitric acid is used, it must be introduced in the form of vapor, or at least as a very fine spray, whereupon it reacts as follows:- 6) 2 S02 + 2 HNO2 + H20 = 2 H2S04 + N203, Perhaps this reaction really occurs in two stages, thus:- (a) S02 + HNO3 = S02.(0H).

(ONO); (b) 2 S02.(OH).(ONO) + H20 =2 H2S04 + N203. The formula assigned to the nitrosylsulphuric acid may perhaps be written: and the compound would then be called nitrosulphonic acid. But in either ease the existence of the substance is only transitory, it being broken up at once by the steam and sulphur dioxide present when the process is working properly. In case there is a deficiency of water vapor in the chambers, and especially if the temperature falls too low, the nitrosylsulphuric acid may separate as crystals, which deposit at various points on the walls, forming "chamber crystals" This is an undesirable accident, for when steam or water come in contact with them they decompose into sulphuric acid, nitric acid, and nitrogen peroxide (N204) :- 4 S02 . (OH) . (ONO) + 2 H20 = 4 H2S04 + N204 + 2 NO.

Then the nitrogen peroxide unites with some of the water, N204 + H20 = HN02 + HNO3, forming nitrous and nitric acids directly on the walls, corroding the lead at the point where the cluster of crystals was attached. To prevent this separation of "chamber crystals" and the retention of nitrogen oxides in the sulphuric acid. an excess of steam in the lead chambers is preferred, although it dilutes the acid somewhat. The raw materials employed in sulphuric acid making are:- 1. Sulphur. (b) Metallic sulphides, such as iron pyrites, chalcopyrite (copper pyrites), sphalerite (zinc blende), etc. (c) Hydrogen sulphide (seldom used). 2. Sodium nitrate or nitric acid. 3. Water as steam. 4. Oxygen as air. The acid may be made from brimstone, pyrites, blende, or hydrogen sulphide, but they are not used together.

Since the burners and chambers employed for one source of sulphur cannot be adapted to that of another without extensive alterations, the manufacturer must decide what material he will use, and erect his plant accordingly. Crude sulphur from the calceroni gives a very pure acid free from arsenic, iron, copper, or zinc, and much smaller condensing chambers may be used for a given yield than when pyrites or blende is employed. The sulphur is placed in iron retorts or on trays in brick ovens, and ignited. The combustion is easily controlled by regulating the amount of air admitted to the retort. Usually the hot vapors from the brimstone burners are passed through a narrow flue or passage, into which a regulated supply of ail' is admitted. This insures complete combustion of any sulphur that may distil owing to too great heat in the retort. To prevent clogging and loss, as much of the sulphur as possible is converted to dioxide. Pyrites, or natural disulphide of iron (FeS2), is a dense, hard mineral of crystalline structure and pale yellow color. The largest deposits in the United States are in Virginia, at Mineral City, and at Charlemont in Massachusetts.

Of the foreign deposits, those in Spain are the most important. A pure pyrites contains 53.3 per cent of sulphur, but that commonly used for acid making carries from 43 to 48 per cent. It seldom pays to use an ore with less than 35 per cent of sulphur, for it will not support its own combustion. The first proposal to use pyrites originated with an Englishman named Hill, who took out a patent for the process in 1818. But it was not until 1838, when the Sicilian government sold the monopoly of the sulphur export to a French firm which nearly trebled the price of crude brimstone, that pyrites began to find favor with acid makers, At the present time, because it is cheap and easily obtained, pyrites has almost completely replaced sulphur for acid making. The product from pyrites is usually contaminated with arsenic , and often with zinc , copper , and selenium . By the oxidation of pyrites in a suitable furnace, the sulphur is converted to dioxide, and iron oxide remains.

The reaction may be written as follows:- 2 FeS2 + 11 0 =4 S02+Fe203 This is not exact, however, as some sulphur remains in the ore and some sulphur trioxide is formed. The proper regulation of the pyrites burners is one of the problems of the manufacturer. If the ore contains over 35 per cent of sulphur, the burning, once started, generates sufficient heat to maintain the combustion, and no fuel is necessary. But zinc sulphide and the" mattes" from metallurgical processes must be heated by fuel. The complete burning of pyrites is difficult. With lump ore there is apt to be a kernel in the centre of the lump, from which the sulphur is not burned out. If the temperature rises too high, the charge fuses together, forming clinkers or "scar," and choking the furnace. If too much air is admitted, the furnace cools below the temperature at which fresh pyrites will ignite, and the gases leaving the burner are so diluted that the desired reactions do not take place in the lead chambers.

With "smalls" the tendency to fuse is more marked than with lump ore, and the fine ore packs together so densely that the air will not penetrate it, and unless it is constantly stirred only the surface is burned. (The lump ore is that which has been broken to about the size of a goose egg, the "smalls" constituting what will pass through a half-inch screen.) Pyrites burners are usually built in benches containing from three to thirty furnaces, in order that the supply of gas may not be broken, while charging or cleaning one furnace. A burner for lump ore (Fig. 21) * consists of a brick furnace, containing a grate formed of single loose iron bars (B, B) having a square section, and resting in grooves at each end. These bars may be turned parallel with their longitudinal axes, but have no lateral motion. They are so adjusted that their sides are at an angle of 45° to the vertical. After a charge is burned, the bars are given several quarter turns by means of a key, to allow the cinders on them to drop through into the ash pit. Ail' is admitted by dampers beneath the grate.

When properly working, the cinders resting on the bars are nearly cold, the hottest part of the fire being eighteen inches above the grate. The furnaces are lined with fire-brick, and to prevent any access of air except through the dampers, the doors (0, D) for charging, cleaning, raking, etc., are make to fit closely, and are generally luted with clay. All the burners in one bench deliver their sulphur dioxide gas into a common, wide flue, or "dust box" (F), where any fine dust carried along by the gases may settle before they enter the Glover towel'. This dust consists of unburned pyrites, arsenic, antimony or zinc oxides, iron oxide, etc. In one or more of the burners a cast-iron "nitre pot" (P) may be set, in which nitrous gases are generated by the action of sulphuric chamber built into the flue (F), and heated by the waste heat of the burners. Sometimes, however, the pots are placed in separate furnaces. A lump burnersof average size has a grate area of 15 to 25 square feet.

The furnace is sometimes made slightly hopper-shapedinside, so that it is larger at the level of the charging doors than at the grate bars.' About 40 pounds of pyrites, containing 48 per cent of sulphur, are burned per square foot of grate area in 24 hours, a larger quantity of such high-grade ore being liable to cause fusion, unless great care is exercised. A larger quantity of poorer ore may be burned daily, without danger of fusion.

Question: Which acid used to be made by the lead chamber process but is now usually made by the contact process?
Answer: Sulphuric